Chemistry: Acid Base and Solution Equilibria

JF Chemistry 1101 2010 Introduction to Physical Chemistry: Acid Base and Solution Equilibria.

Dr Mike Lyons School of Chemistry [email protected]

Required Reading Material.

• Silberberg, Chemistry, 4th edition. – Chapter 18. • Acid/base equilibria. pp.766-813. –Chapter 19. • Ionic equilibria in aqueous systems. pp.814-862. • Kotz, Treichel and Weaver, 7th edition. – Chapter 17 (Chemistry of Acids and Bases) & Chapter18 (Principles of reactivity: other aspects of ionic equilibria), pp.760-859. •Chemistry3, Burrows et al. – Chapter 6, Acids & bases, pp.263-299.

Review : Kotz Chapter 3 for simple acid/base definitions.

Lecture 13.

Acid/base chemistry : Simple ideas: Arrhenius, Bronsted-Lowry, Lewis.

Kotz: section 3.7, pp.131-139. Section 17.1, pp.761-762. Acid and Bases

2 Acid and Bases

Acid and Bases

3 Arrhenius (or Classical) Acid-Base Definition

•An acid is a neutral substance that contains hydrogen and HCl → H + (aq) + Cl − (aq) dissociates or ionizes in water to yield hydrated protons or NaOH → Na+ (aq) + OH − (aq) + hydronium ions H3O . •A base is a neutral substance that contains the hydroxyl group HCl + NaOH → NaCl + H 2O and dissociates in water to yield hydrated hydroxide ions OH -. • Neutralization is the reaction of + + an H (H3O ) ion from the acid and the OH - ion from the base to form water, H2O. • These definitions although correct are limited in that they are not very general and do not • Give a comprehensive idea of what acidity and basicity entails.

+ + Arrhenius acid is a substance that produces H (H3O ) in water.

Arrhenius base is a substance that produces OH- in water.

4 Acids and bases: Bronsted/Lowry definition.

• Bronsted/Lowry Acid (HA): – An acid is a species which donates a proton • Bronsted/Lowry Base (B): – A base is a species which accepts a proton. • These definitions are quite general and refer to the reaction between an acid and a base. • An acid must contain H in its - formula; HNO3 and H2PO4 are two • In the Brønsted-Lowry perspective: examples, all Arrhenius acids are Brønsted-Lowry acids. one species donates a proton and • A base must contain a lone pair of another species accepts it: an acid- electrons to bind the H+ ion; a few base reaction is a proton transfer 2- - examples are NH3, CO3 , F , as well as OH -. Brønsted-Lowry bases are process. not Arrhenius bases, but all 3 Arrhenius bases contain the Chemistry section 6.1. pp.264-267. Brønsted-Lowry base OH-. Kotz 7th ed. Section 17.1. pp.761-765

BL acid/base equilibria. BL base BL acid + - HA(aq) + H2O (l) H3O (aq) + A (aq) Water can function Proton transfer both as an acid BL acid and a base depending on the B (aq) + H O (l) HB+ (aq) +OH- (aq) BL base 2 circumstances.

• Proton donation and acceptance are dynamic processes for all acids and bases. Hence a proton transfer equilibrium is rapidly established Acid : proton donor in solution. Base : proton acceptor • The equilibrium reaction is described in terms of conjugate acid/base pairs. • The conjugate base (CB) of a BL acid is the base which forms when the acid has donated a proton. • The conjugate acid (CA) of a BL base is the acid which forms when the base has accepted a proton. • A conjugate acid has one more proton than the base has, and a conjugate base one less proton than the acid has. • If the acid of a conjugate acid/base pair is strong (good tendency to donate a proton) then the conjugate base will be weak (small tendency to accept a proton) and vice versa. Proton transfer

HA (aq) + B (aq) BH+ (aq) + A- (aq) AB CACB

5 A Brønsted acid is a proton donor A Brønsted base is a proton acceptor

base acid acid base conjugate conjugate base acid acid base

15.1

Brønsted-Lowry Acid-Base Definition

An acid is a proton donor, any species which donates a H+.

A base is a proton acceptor, any species which accepts a H+.

An acid-base reaction can now be viewed from the standpoint of the reactants AND the products. An acid reactant will produce a base product and the two will constitute an acid-base conjugate pair.

6 Table 18.4 The Conjugate Pairs in Some Acid-Base Reactions Conjugate Pair

Acid + Base Base + Acid

Conjugate Pair

- + Reaction 1 HF+ H2O F + H3O

Reaction 2 HCOOH+ CN- HCOO- + HCN

+ 2- - Reaction 3 NH4 + CO3 NH3 + HCO3

- - 2- Reaction 4 H2PO4 + OH HPO4 + H2O

+ - 2+ Reaction 5 H2SO4 + N2H5 HSO4 + N2H6

2- 2- 3- - Reaction 6 HPO4 + SO3 PO4 + HSO3

15.4

7 SAMPLE PROBLEM 18.4: Identifying Conjugate Acid-Base Pairs

PROBLEM: The following reactions are important environmental processes. Identify the conjugate acid-base pairs.

- 2- 2- - (a) H2PO4 (aq) + CO3 (aq) HPO4 (aq) + HCO3 (aq)

2- - - (b) H2O(l) + SO3 (aq) OH (aq) + HSO3 (aq) PLAN: Identify proton donors (acids) and proton acceptors (bases). conjugate pair conjugate pair1 2

- 2- 2- - SOLUTION: (a) H2PO4 (aq) + CO3 (aq) HPO4 (aq) + HCO3 (aq) proton proton proton proton donor acceptor acceptor donor conjugate pair conjugate pair1 2

2- - - (b) H2O(l) + SO3 (aq) OH (aq) + HSO3 (aq) proton proton proton proton donor acceptor acceptor donor

8 Strong and weak acids. Strong acids dissociate completely into ions in water: + - HA(g or l) + H2O(l) H3O (aq) + A (aq) In a dilute solution of a strong acid, almost no HA molecules + exist: [H3O ]][ = [HA ]init or [[]HA]eq = 0 [H O+][A-] 3 at equilibrium, Q = K >> 1 Qc = c c [HA][H2O] + - Nitric acid is an example: HNO3 (l) + H2O(l) H3O (aq) + NO3 (aq) Weak acids dissociate very slightly into ions in water: + - HA(aq) + H2O(aq) H3O (aq) + A (aq) In a dilute solution of a weak acid, the great majority of HA + molecules are undissociated: [H3O ] << [HA]init or [HA]eq = [HA]init + - [H3O ][A ] Qc = at equilibrium, Qc = Kc << 1 [HA][H2O]

Complete ionization The Extent of Dissociation for Strong and Weak Acids

Key concept : Acid/base strength quantified in terms of extent or degree of dissociation. An acid or base is classified as strong if it is fully ionized in solution (e.g. HCl, NaOH). An acid or base is classified as weak if only a small fraction is ionized in solution (e.g. CH COOH, NH ). 3 3 Partial ionization

+ - Strong acid: HA(g or l) + H2O(l) H3O (aq) + A (aq)

Extent of dissociation: weak acid.

+ - Weak acid: HA(aq) + H2O(l) H3O (aq) + A (aq)

10 Reactivity of strong and weak acids.

1M HCl(aq) 1M CH3COOH(aq)

Classifying the Relative Strengths of Acids. Strong acids. There are two types of strong acids:

•The hydrohalic acids HCl, HBr, and HI •Oxoacids in which the number offf O atoms exceeds the number of

ionizable H atoms by two or more, such as HNO3, H2SO4, HClO4

Weak acids. There are many more weak acids than strong ones. Four types, with examples, are:

• The hydrohalic acid HF • Those acid s in w hic h H is boun de d to O or to ha logen, suc h as

HCN and H2S • Oxoacids in which the number of O atoms equals or exceeds by one

the number of ionizable H atoms, such as HClO, HNO2, and H3PO4 • Organic acids (general formula RCOOH), such as CH3COOH and C6H5COOH.

11 Classifying the Relative Strengths of Bases.

Strong bases. • Soluble compounds containing O2- or OH- ions are strong bases. The cations are usually those of the most active metals:

M2O or MOH, where M= Group 1A(1) metals (Li, Na, K, Rb, Cs). • MO or M(OH)2, wh ere M = Group 2A(2) meta ls (Ca, Sr, Ba ) [MgO and Mg(OH)2 are only slightly soluble, but the soluble portion dissociates completely.]

Weak bases. • Many compounds with an electron-rich nitrogen are weak bases (none are Arrhenius bases). The common structural feature is an N atom that has a lone electron pair in its Lewis structure.

• Ammonia (NH3) • Amines (general formula RNH2, R2NH, R3N), such as CH3CH2NH2, (CH3)2NH, (C3H7)3N, and C5H5N

Strong Electrolyte – 100% dissociation

H O NaCl (s)Na2 + (aq) + Cl- (aq)

Weak Electrolyte – not completely dissociated

- + CH3COOH CH3COO (aq) + H (aq)

Strong Acids are strong electrolytes

+ - HCl (aq) +H+ H2O (l) H3O (aq) +Cl+ Cl (aq) + - HNO3 (aq) + H2O (l) H3O (aq) + NO3 (aq) + - HClO4 (aq) + H2O (l) H3O (aq) + ClO4 (aq) + - H2SO4 (aq) + H2O (l) H3O (aq) + HSO4 (aq) 15.4

12 Weak Acids are weak electrolytes

+ - HF (aq) + H2O (l) H3O (aq) + F (aq) + - HNO2 (aq) + H2O (l) H3O (aq) + NO2 (aq) - + 2- HSO4 (aq) + H2O (l) H3O (aq) + SO4 (aq) + - H2O (l) + H2O (l) H3O (aq) + OH (aq)

Strong Bases are strong electrolytes H O NaOH (s) 2 Na+ (aq) + OH- (aq) H O KOH (s) 2 K+ (aq) + OH- (aq) H O 2 2+ - Ba(OH)2 (s) Ba (aq) + 2OH (aq)

15.4

Weak Bases are weak electrolytes

- - F (aq) + H2O (l) OH (aq) + HF (aq) - - NO2 (aq) + H2O (l) OH (aq) + HNO2 (aq)

Conjugate acid-base pairs: • The conjugate base of a strong acid has no measurable strength.

+ •H3O is the strongest acid that can exist in aqueous solution. •The OH- ion is the strongest base that can exist in aqeous solution.

15.4

13 Representing Protons

• Both representations of the proton are equivalent.

+ + + •H5O2 (aq), H7O3 (aq), H9O4 (aq) have been observed. • We will use H+()!(aq)!

+ H (H 2O)3 + H 7O3

The hydrated proton is quite a complex entity. It is usually represented in shorthand form as H+(aq). A better representation is in terms of the + hydronium ion H3O . We will adopt this representation a lot. The real situation + is more complex. The H3O ion binds to other water molecules forming + a mixture of species with the general formula H(H2O)n . In fact the structural details of liquid water is still a hot item of research.

14 What is H+ (aq)?

H + H O+ O+ 3 H H H H+

H 2O OH2 H H + H H O O O H + H H H

H O + 2 H5O2 + H9O4

15 Yet more sophistication: Lewis acidity

+ + An Arrhenius acid is defined as a substance that produces H (H3O ) in water. A Brønsted acid is defined as a proton donor

A LiLewis acid is defin e d a s a subst an ce tha t ca n accept a pai r of electrons. A Lewis base is defined as a substance that can donate a pair of electrons •• + • - •• G.N. Lewis 1875-1946 H + OH• H O H •• ••

acid base See Kotz section 17.9 pp.789-798. H H

+ •

H + N• H H N H H H acid base http://en.wikipedia.org/wiki/Gilbert_N._Lewis

Electron-Pair Donation and the Lewis Acid-Base Definition

The Lewis acid-base definition :

A base is any species that donates an electron pair. An acid is any ssipecies ththtat acceptts an elltectron pai r. Protons act as Lewis acids in that they accept an electron pair in all reactions: B + H+ B H+

The product of any Lewis acid-base reaction is called an adduct, a single species that contains a new covalent bbdond.

A Lewis base has a lone pair of electrons to donate.

A Lewis acid has a vacant orbital

16 Lewis Acid/Base Reaction

Lewis Acids and Bases

F H F H •

F B + NHN• H F B NHN H F H F H acid base

No protons donated or accepted!

15.12

17 Lewis Acids & Bases

Other good examples involve metal ions. •• •O—H •• Co2+ • O—H Co2+ • H ACID BASE H

Lewis Acids & Bases

The combination of metal ions (Lewis acids)

with Lewis bases such as H2O and NH3 leads to COMPLEX IONS

18 2+ Reaction of NH3 with Cu (aq)

PLAY MOVIE

The Lewis AcidAcid--BaseBase Chemistry of Nickel(II)

19 Lewis Acids & Bases

2+ 2+ [Ni(H2O)6] + 6 NH3 Æ [Ni(NH3)6]

+ DMG

DMG = dimethylglyoxime, a standard reagent to detect nickel(II)

Lewis AcidAcid--BaseBase Interactions in Biology

• The heme group in hemoglobin can

interact with O2 and PLAY MOVIE CO. • The Fe ion in hemoglobin is a Lewis acid

•O2 and CO can act as Lewis bases

Heme group

20 Lewis Acids & Bases

Many complex ions containing water undergo HYDROLYSIS to give acidic solutions. 2+ + + [Cu(H2O)4] + H2O Æ [Cu(H2O)3(OH)] + H3O

Lewis Acids & Bases

This explains why water solutions of Fe3+, Al3+, Cu2+, Pb2+, etc. are acidic.

This interaction weakens this bond

Another H2O pulls this H away as H+

21 Amphoterism of Al(OH)3

Lewis Acids & Bases

This explains AMPHOTERIC nature of some metal hydroxides. + 3+ Al(OH)3(s) + 3 H3O Æ Al + 6 H2O

Here Al(OH)3 is a Brønsted base. - - Al(OH)3(s) + OH Æ Al(OH)4

Here Al(OH)3 is a Lewis acid.

•• •• Al3+ O—H- ••